Chemical Bonding and StructureIB CHEMISTRY

TOPIC 4

Bond - A force that holds atoms together and

makes them function as a unit.

4.1 Ionic BondingAn ion is a charged particle. Ions form from atoms or from groups of atoms by loss or gain of one or more electron. The number of electrons lost/gained is determined by the electron configuration. The ideal configuration for any atom is one that is isoelectronic to the noble gases.

The number of charges on the ion formed is equal to the number of electrons lost or gained.

4.1 Ionic Bonding

When a metal atom loses electrons it forms a positive ion, also called a cation.When a non-metal atom gains electrons it forms a negative ion, also called a anion.

4.1 Ionic Bonding When metals lose electrons:

Group 1 = cation with a 1+ charge

Group 2 = cation with a 2+ charge

Group 3 = cation with a 3+ charge

4.1 Ionic Bonding When non-metals lose electrons:

Group 5 = anion with a 3- charge

Group 6 = anion with a 2- charge

Group 7 = anion with a 1- charge

4.1 Ionic Bonding

The attraction between the anion and cation is called the electrostatic attraction.

This attraction is what call an ionic bond.

4.1 Ionic Bonding Under normal conditions, ionic compounds are typically solids with lattice-type structures (3D repeating units of cations and anions)

The ionic bond (electrostatic attraction) is not only between two atoms – it is throughout the lattice!

Think/Pair/Share: How do we determine whether two elements will form an ionic bond?

4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values.

Electronegativity:The measure of the ability of an atom to attract an electron

Difference in electronegativities of the two elements:If greater than or equal to 1.8 units on the Pauling scale, then an ionic bond will be formed.

Pg. 8 of Data Booklet

4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values.

One element will usually be a metal on the left of the Periodic Table and the other a non-metal on the right.

Remember: Tendencies to form cations increases down the groups on the left.

Tendencies to form anions increase up the groups on the right.

State the charge and formula for these:

A) Nitrate

B) Hydroxide

C) Hydrogen carbonate

D) Carbonate

E) Sulfate

F) Phosphate

G) Ammonium

4.2 Covalent Bonding “Share electrons” in order to achieve isoelectronic configuration as noble gas

Electrostatic attraction between positively charge nuclei and a pair of electrons

Small difference in electronegativity of elements that will form covalent bonds

Octet Rule – eight valence electrons (s +p) ◦ Exceptions to the octet rule?

4.2 Covalent BondingWhat if the octet can’t be satisfied with sharing of two electrons?

Multiple electrons can be shared between two atoms

Sharing 1 pair of electrons = single bondSharing 2 pairs of electrons = double bondSharing 3 pairs of electrons = triple bond

4.2 Covalent Bonding •Bond strength (enthalpy) – energy needed to break bond

◦ •Increases as more electrons shared (triple > double > single) ◦ •Greater attraction between nucleus and electrons

•Bond length – distance between two bonded nuclei ◦ •Decreases as more electrons shared (more attraction means closer

together) ◦ •Triple < double < single

Pg. 10 and 11 of Data Booklet

Bond Length (pm) Strength (kJ mol-1)C-C 154 346C=C 134 614C≡C 120 839

4.2 Covalent Bond PolarityElectronegativity -- describes the unequal sharing of bonding electron pairs in a compound. (Larger value = greater pull of one atom over the other, and vice versa.)

Polarity – separation of charges across a molecule

All covalent

bonds are

polar!!

4.2 Covalent Bond Polarity Hydrogen has a partial positive charge (electronegativity of 2.2) Fluorine has a partial negative charge (electronegativity of 4.0)

4.2 Covalent Bond PolarityCompare electronegativity of each atom:If greater than 0.5, one end of the bond will have a larger electron density resulting in partial charges on overall compoundIf less than 0.5, compound is considered non-polar

PracticeArrange the following bonds in order of increasing polarity…

H-H F-H Cl-H S-H O-H

Determine the polarity of the covalent bonds between (need DB): A) nitrogen and fluorine B) hydrogen and carbon C) sulfur and oxygen D) chlorine and chlorine E) oxygen and nitrogen

4.3 Covalent structuresLewis structures show valence electrons of atoms (those in outer most energy level) in covalently bonded species. •Write the symbol of the element surrounded dots representing the amount of valence electrons of the atom (remember Hund’s rule)

•Structures show how atoms come together to achieve stable electron configurations

Na C O [ Cl ]-

4.3 Covalent structuresDraw Lewis Structures for the following:

Boron Fluorine Silicon Krypton Iodide

4.3 Covalent structures When covalent compounds are drawn, show both bonding and non-bonding electrons (lone pairs)

All atoms will follow the octet rule or the “duet rule” (exceptions?)

4.3 Covalent structures

Draw the Lewis structures for the following:

1. Determine total number of e- wanted 2. Determine total number of e- you have 3. Subtract wanted – have and divide by 2 4. This shows the number of pairs shared (either in single, double, or triple bonds).

5. Distribute remaining e- as lone pairs around the atoms.

6. Ions – add 1e- for each neg. charge or subtract 1e- for each positive charge; draw brackets around structure with charge outside

C2H6

C2H4

[OH]-

N2

C6H6

4.3 Covalent structures Resonance structures occur when there is more than one possible location for a double bond in a molecule.

4.3 Covalent structures Draw the resonance structures for:

CO32- O3

Richard Thornely

Special case: coordinate bond Typically, a covalent bond is formed by each atom donating one electron each to the bond. Sometimes, however, an atom can donate a lone pair to form a covalent bond. Chemically, there is no difference between a coordinate bond (or dative covalent bond) and a “regular” covalent bond.

Ex: NH3 + BF3

Special case: coordinate bond AlCl3 is electron deficient (just like BF3), but in the vapor at sublimation temperatures has the formula of Al2Cl6. It exists as a dimer (two molecules joined together).

Draw the Lewis structures for:

a)HF b)CF3Cl c)C2H6 d)NO3

- e)SO2

f)C2H4 g)C2H2 h)NO2

1- i)SO3

2- j)BF3

4.3 Covalent Structure You can find the shape of a molecular structure by knowing where bonds and lone pairs are around a molecule.

Think of it as a 3-D Lewis Structure Valence shell electron pair repulsion theory (VSEPR) – electron pairs around central atom will repel each other to maximize distances between e- pairs (neg. repels neg.)

Results in 3-D shape

What is the total number of atoms in 0.50mol of 1,4-diaminobenzene, H2NC6H4NH2? (Avogadro’s constant (L or NA) = 6.0 ×1023mol–1.)

A. 16.0 x 1023

B. 48.0 x 1023

C. 96.0 x 1023

D. 192.0 x 1023

4.3 Covalent Structure - VSEPR 1. Count number of atoms bonded to central atom (single, double or triple bonds count as 1 bond)

2. Count number of lone pairs around central atom (each lone pair is 1 bond)

3. Total number of bonds and lone pairs (for the central atom) determines overall shape

Results:

2 = Linear (180˚)

3 = Planar Triangular (120˚)

4 = Tetrahedral (109.5˚)

4.3 Covalent Structure - VSEPR Lone pairs cause greater repulsion in tetrahedral arrangements Alters the angles/overall appearance

4.3 Covalent Structure - VSEPRDraw Lewis Structures, determine bond angles, and describe the shape for the following molecules…

BF3

NO3-

H2O

SO2 (watch out for lone pairs!)

4.3 Polarity of Molecules Not all molecules with polar bonds are polar themselves!

Must take into account shape of molecule

If bonds of equal polarity are arranged symmetrically the polarity is neutralized (non-polar)

Ex. Carbon dioxide – no dipole

4.3 Polarity of MoleculesWater however is still polar. Why?

4.3 Polarity of Molecules Carbon tetrachloride - polar or non-polar?

4.3 Polarity of Molecules Ammonia - polar or non-polar?

Crystalline Solids Some molecules are able to form crystalline structures similar to ionic compounds except they are covalently bonded

Referred to as giant molecular structure or macromolecule Allotropes (different crystalline forms based on bonding)

Carbon has 3 different allotropes.

Allotropes of carbon (a) Diamond – each carbon is bonded to 4 other carbon atoms (tetrahedral structure)

-- each bond equally strong (no plane of weakness)

Allotropes of carbon(b) Graphite – each carbon atom bonded to 3 others in planes

-- each carbon has pair of delocalized, non-bonded electrons (conductor) -- planes attracted by weak IMFs

(c) Graphene – single plane of graphite

Allotropes of carbon (d) Buckminsterfullerene (C60) – each carbon atom is bonded to 3 others in rings that form a sphere

SiO2 – giant covalent Same group as Carbon – similar properties

•Tetrahedral structure (like diamond)

•Silicon dioxide (quartz) – silicon forms single covalent bonds with oxygen

•Each silicon is surrounded by 4 oxygen (SiO2 is the ratio)

•Insoluble in water, high melting point, non-conductor

Crash Course Chemistry #35: The Internet’s Favorite Element – Silicon (9:26)

4.4 Intermolecular Forces • Forces that exist between molecules • Weaker than intramolecular bonds • Helps explain physical properties – solubility, volatility, melting/boiling points • Three different types of IMFs • Van der Waals’ Forces (London Dispersion Forces) • Dipole-Dipole Forces • Hydrogen Bonding

4.4 IMFs •London (dispersion) forces– temporary dipoles that occur when electrons congregate in one area

•Can induce another dipole in neighboring molecules

•Increase with increasing mass •Explains attraction in noble gases and non-polar molecules (diatomic elements for example)

IMFs

4.4 IMFs Dipole-Dipole Forces – attraction between polar molecules

-- permanent dipoles exist with molecules

Forces increase with increasing degree of polarity

Stronger than Van der Waals’ Forces

4.4 IMFsHydrogen Bonding – dipole-dipole attraction (strongest IMF)

Results when hydrogen is bonded to strong electronegative atom (F, O, N)

Hydrogen has strong positive charge (no shielding)

4.4 IMFsWater can form up to four hydrogen bonds with neighboring molecules (because of two H atoms and O’s two lone pairs of electrons)

-- Frozen water takes on a tetrahedral shape -- Water expands upon taking on a solid state -- Proven by smaller density (why ice floats)

4.5 Metallic Bonding Electrostatic attraction between a lattice of positive ions and delocalized electrons

Delocalized electrons – easily detached from nuclei and able to move about freely

Malleability – can slide over each other without breaking more bonds than are made

Conductivity – electrons can move through the structure

Impurities make metals less malleable and ductile

Alloys – mixture of metals will increase strength